Which of the following pairs represents isoelectronic ions?

1. Atomic Number and Subatomic Particles (basic)

Welcome to your first step in mastering chemistry! To understand how the world around us is built, we must look at the Atomic Number—the fundamental "identity card" of an element. The atomic number (represented by the symbol Z) is defined strictly by the number of protons found in the nucleus of an atom. For example, any atom with exactly 6 protons is Carbon, and any atom with 11 protons is Sodium Science, Carbon and its Compounds, p.59. While other particles might change, the number of protons remains constant for a specific element.

An atom is composed of three primary subatomic particles, each carrying a specific charge and role:

• Protons: Positively charged (+) and located in the nucleus. They determine the element's identity.
• Neutrons: Carrying no charge (neutral), they sit in the nucleus with protons and contribute to the atom's mass.
• Electrons: Negatively charged (-) and found orbiting the nucleus in various shells.

In a neutral atom, the number of protons and electrons is perfectly balanced. However, atoms often seek stability by losing or gaining electrons. When a neutral atom like Sodium (Atomic Number 11) loses an electron, it still has 11 protons but now has only 10 electrons. This creates a net positive charge, turning the atom into a cation (Na⁺) Science, Metals and Non-metals, p.46. Conversely, gaining an electron creates an anion with a negative charge. This movement of electrons is the "currency" of chemical reactions, yet the number of atoms of each element remains the same before and after a reaction, adhering to the Law of Conservation of Mass Science, Chemical Reactions and Equations, p.3.

Sources: Science, Carbon and its Compounds, p.59; Science, Metals and Non-metals, p.46; Science, Chemical Reactions and Equations, p.3

2. Electronic Configuration and Shells (basic)

To understand how matter behaves, we must look at how electrons are organized within an atom. This arrangement is called the electronic configuration. Electrons do not just swarm randomly; they reside in specific energy levels or shells, labeled K, L, M, and N, starting from the one closest to the nucleus. Each shell has a maximum capacity: the K shell can hold 2 electrons, while the L and M shells generally aim for 8 to achieve a stable state known as an octet. As we see in Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.47, elements like Helium (2) and Neon (2, 8) have completely filled shells, making them "Noble" or chemically unreactive.

Most elements, however, have incomplete outer shells. The electrons in this outermost shell are called valence electrons. The chemical reactivity of an element is essentially its "hunger" to reach a stable noble gas configuration. Atoms achieve this by either losing, gaining, or sharing electrons. For instance, Sodium (Na) has an atomic number of 11 with a configuration of (2, 8, 1). It is much easier for Sodium to lose that 1 lonely electron in its M shell than to find 7 more. When it loses that electron, its L shell (which has 8 electrons) becomes the new outermost shell, creating a stable Sodium cation (Na⁺) Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.46.

Conversely, non-metals like Chlorine (atomic number 17, configuration 2, 8, 7) are just one electron short of an octet. They tend to gain an electron to form an anion (Cl⁻). This drive to attain a full outer shell is the fundamental reason behind the formation of chemical bonds Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.59. When different atoms or ions end up with the same total number of electrons and the same configuration, we call them isoelectronic species.

Common Electronic Configurations:

Element	Atomic Number	Shell Distribution (K, L, M)	Nature
Helium (He)	2	2	Stable Noble Gas
Magnesium (Mg)	12	2, 8, 2	Reactive Metal
Oxygen (O)	8	2, 6	Reactive Non-metal
Argon (Ar)	18	2, 8, 8	Stable Noble Gas

Sources: Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.46-47; Science, Class X (NCERT 2025 ed.), Carbon and its Compounds, p.59-60

3. Valency and the Octet Rule (basic)

At the heart of chemistry lies a simple quest: the search for stability. Every atom 'desires' to reach a state of minimum energy, which it achieves by having a completely filled outermost shell. This principle is known as the Octet Rule. Much like how we seek financial or professional stability, atoms look at Noble Gases—like Neon (Ne) or Argon (Ar)—as the gold standard because their valence shells are already full Science, class X (NCERT 2025 ed.), Chapter 3, p.47. Except for Helium, which is stable with just 2 electrons (a 'duplet'), most atoms strive to have 8 electrons in their outermost shell to become chemically inert.

The Valency of an atom is essentially its 'combining capacity'—the number of electrons it must lose, gain, or share to achieve that stable octet. For example, a Sodium (Na) atom has 11 electrons, with a configuration of 2, 8, 1. It is far easier for Sodium to lose that 1 lonely electron than to hunt for 7 more. When it loses that electron, its L-shell becomes the new outermost shell, now boasting a perfect 8 Science, class X (NCERT 2025 ed.), Chapter 3, p.46. Conversely, Chlorine (Cl) has 7 electrons in its outer shell; its valency is 1 because it only needs to gain or share 1 electron to complete its octet Science, class X (NCERT 2025 ed.), Chapter 4, p.60.

Element	Atomic Number	Electronic Configuration	Valency
Magnesium (Mg)	12	2, 8, 2	2 (loses 2)
Oxygen (O)	8	2, 6	2 (gains 2)
Nitrogen (N)	7	2, 5	3 (shares 3)

Whether through ionic bonding (transferring electrons) or covalent bonding (sharing electrons, as seen in N₂ or H₂O molecules), the ultimate goal remains the same: attaining the electronic configuration of the nearest noble gas Science, class X (NCERT 2025 ed.), Chapter 4, p.60.

Sources: Science, class X (NCERT 2025 ed.), Chapter 3: Metals and Non-metals, p.46-47; Science, class X (NCERT 2025 ed.), Chapter 4: Carbon and its Compounds, p.60

4. Formation of Ions: Cations and Anions (intermediate)

In nature, atoms are like people—they generally prefer to be in a state of maximum stability. For most atoms, this stability is achieved when they have a completely filled outer electron shell, similar to the noble gases (like Neon or Argon). To reach this "happy" state, atoms often engage in a process called ionization, where they either lose or gain electrons. Since electrons carry a negative charge, any change in their number disrupts the electrical balance of the atom, turning a neutral atom into a charged species called an ion.

Cations are positively charged ions. They are typically formed by metals, which have a few "extra" electrons in their outermost shell. Rather than trying to find many more electrons to fill the shell, it is energetically easier for them to simply give away those few outer electrons. For example, a Magnesium (Mg) atom has 12 electrons. By losing 2 electrons, it becomes a Mg²⁺ cation, achieving the stable electronic configuration of Neon. As noted in Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.48, these cations represent the "positive" half of many chemical compounds.

Anions, on the other hand, are negatively charged ions. These are usually formed by non-metals, which are just a few electrons short of a full shell. Instead of giving up many electrons, they prefer to "steal" or gain them. When a Chlorine (Cl) atom gains one electron, it becomes a Cl⁻ anion. The result of this "give and take" is a powerful electrostatic attraction between the newly formed positive and negative ions, leading to the creation of ionic compounds like Magnesium Chloride (MgCl₂). This fundamental transfer of electrons is the bedrock of chemical reactivity between metals and non-metals Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.56.

Feature	Cation	Anion
Charge	Positive (+)	Negative (–)
Formation	Loss of electrons	Gain of electrons
Type of Element	Mostly Metals	Mostly Non-metals
Example	Na⁺, Mg²⁺, Al³⁺	Cl⁻, O²⁻, S²⁻

Sources: Science, class X (NCERT 2025 ed.), Chapter 3: Metals and Non-metals, p.48; Science, class X (NCERT 2025 ed.), Chapter 3: Metals and Non-metals, p.56

5. Periodic Trends: Atomic and Ionic Radii (intermediate)

To understand the geometry of the microscopic world, we must look at Atomic and Ionic Radii. Atomic radius is essentially the distance from the center of the nucleus to the outermost shell of electrons. As we move across a period (left to right) in the periodic table, the atomic radius decreases. This happens because the number of protons (nuclear charge) increases, pulling the electrons closer to the nucleus. Conversely, as we move down a group, the radius increases because new electron shells are added, making the atom larger. When atoms gain or lose electrons to become stable, they form ions. A cation is a positively charged ion formed when an atom loses electrons, while an anion is a negatively charged ion formed when an atom gains electrons Physical Geography by PMF IAS, Thunderstorm, p.348. A cation is always smaller than its parent atom. For instance, when a Sodium (Na) atom loses an electron to become Na⁺, its entire outer shell (the M shell) disappears, and the remaining 10 electrons are pulled even more tightly by the 11 protons in the nucleus Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.46. In contrast, an anion is always larger than its parent atom because the addition of electrons increases electron-electron repulsion, pushing them further apart.

Ion Type	Charge	Formation	Size Change
Cation	Positive (+)	Loss of electrons	Smaller than parent atom
Anion	Negative (-)	Gain of electrons	Larger than parent atom

A fascinating subset of this topic involves isoelectronic species—atoms or ions that have the same number of electrons. For example, O²⁻, F⁻, Na⁺, and Mg²⁺ all possess 10 electrons. However, they do not have the same size! Their radii differ based on their nuclear charge (number of protons). The species with the highest number of protons will pull the electrons most strongly, resulting in the smallest radius. In our example, Mg²⁺ (12 protons) would be the smallest, while O²⁻ (8 protons) would be the largest.

Sources: Science, class X (NCERT 2025 ed.), Metals and Non-metals, p.46; Physical Geography by PMF IAS, Thunderstorm, p.348

6. Understanding Isoelectronic Species (exam-level)

In chemistry, the prefix 'iso-' translates to 'same'. When we talk about isoelectronic species, we are referring to atoms, ions, or molecules that share the exact same number of electrons and, consequently, the same electronic configuration. This is a fundamental concept because atoms often gain or lose electrons to achieve the stable electronic arrangement of the nearest noble gas, a process driven by the desire for chemical stability. Science, Class X (NCERT 2025 ed.), Chapter 3, p. 47.

To identify whether two species are isoelectronic, you must calculate their total electron count. For a neutral atom, the number of electrons equals its atomic number (Z). However, ions are formed when atoms shift away from this neutrality:

• Cations: Positively charged ions formed by losing electrons. (Number of electrons = Z - Charge).
• Anions: Negatively charged ions formed by gaining electrons. (Number of electrons = Z + |Charge|). Physical Geography by PMF IAS, Thunderstorm, p. 348.

Even though these species have the same number of electrons, they are not identical; they differ in their nuclear charge (number of protons), which affects their size and reactivity.

Let's look at the Argon (Ar) family as a classic example. Argon has an atomic number of 18, meaning it has 18 electrons. Other elements near it in the periodic table will strive to reach this 18-electron count through ionization:

Species	Atomic Number (Protons)	Change	Total Electrons
Argon (Ar)	18	None (Neutral)	18
Potassium Ion (K⁺)	19	Lost 1 electron	18
Calcium Ion (Ca²⁺)	20	Lost 2 electrons	18
Chloride Ion (Cl⁻)	17	Gained 1 electron	18
Sulfide Ion (S²⁻)	16	Gained 2 electrons	18

Sources: Science, Class X (NCERT 2025 ed.), Chapter 3: Metals and Non-metals, p.47; Physical Geography by PMF IAS, Thunderstorm, p.348

7. Solving the Original PYQ (exam-level)

Now that you have mastered the fundamentals of atomic structure and the process of ion formation, this question serves as the perfect synthesis of those concepts. To solve this, you must apply the definition of isoelectronic species: atoms or ions that possess the exact same number of electrons. This requires you to bridge your knowledge of the periodic table with the mathematical adjustment of electron counts based on positive or negative charges.

Let’s walk through the reasoning as if we were in the exam hall. Focus first on the atomic number, which represents the neutral state. In the correct option, (D) Ca²⁺, S^2-, we see this balance in action. Calcium (atomic number 20) loses two electrons to form a cation, resulting in 18 electrons. Sulfur (atomic number 16) gains two electrons to form an anion, also resulting in 18 electrons. Since both reach the stable configuration of the noble gas Argon, they are isoelectronic. As noted in Science, class X (NCERT 2025 ed.), this state is the driving force behind most chemical reactions.

It is crucial to recognize the common UPSC traps present in the other options. Options (A) and (C) pair elements from the same vertical group (Na/K and Mg/Ca). While these ions share the same valency and similar chemical properties, they belong to different periods, meaning their total electron counts are different. Do not confuse "similar charge" with being "isoelectronic." Always perform the subtraction for cations and addition for anions to ensure the final counts match perfectly before selecting your answer.